Electrocatalytic activity of electropolymerized methylene blue … · 2019. 5. 10. · oxidation...

Electrocatalytic Activity ofElectropolymerized Meldola’s Blue towardOxidation of Dopamine

著者 Yamaguchi Takahiro, Komura Teruhisa, HayashiSayomi, Asano Manabu, Niu Guang Yao, TaahashiKohshin

journal orpublication title

Electrochemistry = 電気化学および工業物理化学

volume 74number 1page range 32-41year 2006-01-01URL http://hdl.handle.net/2297/36482

doi: 10.5796/electrochemistry.74.32

Electrocatalytic Activity of Electropolymerized Meldola’s blue toward Oxidation of Dopamine

Takahiro YAMAGUCHI*, Teruhisa KOMURA, Sayomi HAYASHI, Manabu ASANO,

Guang Yao NIU, Kohshin TAKAHASHI

Division of Material Sciences, Graduate School of Natural Science and Technology,

Kanazawa University (Kakumamachi, Kanazawa, 920-1192, Japan)

1

To understand the kinetics of charge transfer from chemically modified electrodes to solution

species, the authors have investigated the influence of poly(Meldola’s blue) on the kinetic parameters of

dopamine oxidation by rotating disc electrode voltammetry. The polymer film fixed on electrode

surfaces raised the dopamine peak current by a factor of 10 in a 0.1 mM solution; this electrocatalytic

effect enables one to detect dopamine at concentrations of the order of 5 µM. The polymer film

increased the standard heterogeneous rate constant of dopamine oxidation more than 10-fold and favored

a two-electron transfer step, thus raising the dependence of the kinetic current on the electrode potential.

These effects were independent of electron self-exchange between redox-active sites in the film. They

arose from the incorporation of dopamine with the polymer. This binding interaction decreased with

protonation of the polymer, because of electrostatic repulsion between the positively charged species.

Keywords: Chemically modified electrode, Electrocatalytic oxidation, Poly(Meldola’s blue), Dopamine,

Rotating disc electrode

1. Introduction

Dopamine (3,4-dihydroxyphenethylamine, DA) is a catechol-based neurotransmitter which plays

an important role in the mammalian central nervous system. Because DA is rather irreversibly oxidized

at metal electrodes, its redox reaction requires a high overpotential. In addition, DA is present at a very

low concentration of 0.01−1 µM in the extra-cellular fluid of the central nervous system and has an

oxidation potential close to that of ascorbic acid coexisting at a higher concentration1−4). Appropriate

electrocatalysts are hence necessary for electrochemical detection of DA5−8). A number of chemically

modified electrodes have been proposed for the determination of this compound; metal electrodes coated

with thin films such as metal(II)hexacyanoferrate9−11), metal phthalocyanines12,13), and polymers14−17).

Several workers have reported sensitive or selective voltammetric monitoring of DA1, 8−20). Miller and

Zhou21) prepared a composite poly(N-methylpyrrole)-polystyrenesulfonate film to incorporate cationic

DA into the reduced polymer. Piro and others22) have recently developed a new cation-exchanging

polymer, poly(5-amino-1,4-naphtoquinone), which can uptake DA up to 100 nmol cm-2 in a 300 nm

thick film. The DA incorporated was released by applying the oxidative potential to the redox-active

polymer. Electrocontrolled release of bio-related materials by polymer-modified electrodes plays a

crucial role in delivering drugs at a constant rate to a precise site.

Azines (phenazines, phenothiazines, and phenoxazines) undergo a reversible two-electron,

one-proton reduction to yield the corresponding 10-hydro-azine forms; thus these dyes have been often

2

used as redox indicators and mediators in the field of biotechnological analysis23−27). In addition, most of

the amino group-substituted azines are irreversibly oxidized to their respective radical cations in aqueous

media and have yielded polymeric films on electrode surfaces27−29), though showing a low current

efficiency. Having redox-active properties like those of the corresponding monomers, the

electropolymerized azines have attracted a growing interest because of their analytical application23, 30−33).

Their current-enhancing effects on the hydrogenation-dehydrogenation of myoglobin34), hemoglobin34),

and NADH28, 35, 36) prompted us to examine the electrocatalytic activity of polyazines against oxidation of

DA.

Voltammetric detection of solution species and electrocatalytic efficiency at polymer-modified

electrodes depend primarily on how greatly a surface-immobilized polymer enhances the rate of the

redox reaction of a target. Hence, the following subjects need to be investigated: (1) the kinetics of

oxidation-reduction of the polymer and (2) its catalytic activity toward the target redox reaction. To

understand the mediating properties of polyazines, one requires to examine the rate of coupled electron

and proton transport in the electroactive nitrogen-including polymers, because the kinetics of their redox

reactions depend on both the reduction and protonation levels. The preceding paper37) demonstrates that

the diffusion coefficient of charge carriers in a polyazine film increases with decreasing solution pH.

This change results from electron hopping accompanied by proton transfer. On the other hand, limited

data on electron exchange between the catalytic center and solution substrate have hampered a thorough

understanding of the selective electrocatalysis due to polyazines.

This paper describes the influence of poly(Meldola’s blue) (PMD) films on the kinetic parameters

of DA electrooxidation and ascribes the enhanced oxidation rate to the incorporation of DA into the

polymer, which arises from their binding interaction.

2. Experimental

Meldola’s blue MD+ (Structure 1), methylene blue, and DA hydrochloride were obtained from

Kanto Kagaku Co., and used as received. Sodium sulfate (Wako Chemicals, purity 99.9%) was

recrystallized from doubly distilled water and then used as a supporting electrolyte. Thin films of PMD

were electrolytically deposited on glassy carbon (GC) of 0.071 cm2 area by the same method as that

used for poly(methylene blue) 37). The potential of the working electrode was repetitively cycled

between -0.4 and 1.6 V (versus Ag/AgCl) at a scan rate of 100 mV s-1 in a 1 mM MD+ solution (pH 8)

containing 0.2 M Na2SO4. Sample solution contained 20 mM buffer salts, either acetate(pH 4−6),

phosphate(pH 6−8) depending on the required pH.

3

The surface coverage Θ of the resulting polymer, the surface concentration of a two-electron

reaction site, was calculated from the voltammetric charge (Q) consumed in reducing the entire polymer

at a scan rate of 5 mV s-1 in a pure supporting electrolyte solution (pH 4−5):

Θ = Q /2FS (1)

where F is the Faraday constant and S is the electrode area. The polymer film was too thin (less than 50

nm) to estimate the film thickness by scanning electron microscopy. The thickness of PMD films given

below was calculated from the Θ versus thickness plot measured for poly(methylene blue) 37), assuming

that the two polymers are equal in the concentration of electroactive sites. The attenuated total reflection

(ATR) infrared spectra of PMD films deposited on electrodes were recorded on a Horiba FT-710 IR

spectrophotometer. In situ visible absorption spectrum of the polymer was measured at its different

reduction levels with a 10 mm path-length quartz cell containing a polymer-coated indium tin oxide

electrode, a platinum wire, and an Ag/AgCl reference electrode.

All the electrochemical experiments were carried out at a temperature of 20 ± 5 ºC under a

nitrogen atmosphere with a conventional three-electrode glass cell enclosed in a grounded Faraday cage.

A 3 mm diameter GC rod mounted with epoxy resin into a glass tube was used as a working electrode, a

platinum gauze of 20 cm2 area as a counter electrode, and an Ag|AgCl|3.4 M KCl electrode as the

reference electrode. The working electrode was polished with aqueous slurries of alumina (down to a

grain size of 0.05 µm) and then sonicated in purified water for 5 min. Cyclic voltammetry, differential

pulse voltammetry, and potential step chronocoulometry were performed by using a Hokuto Denko

HZ-3000 electrochemical system consisting of a FMV Desk Power CIX 407c personal computer, a

HAG-1510m potentiostat, and a BJC-F 200 printer. A Hokuto Denko rotating disk electrode system

with a glassy carbon electrode (diameter 0.5 cm) was used in the rotating disk electrode experiment.

3. Results and discussion

3.1. Redox properties of PMD

When oxidized at 1.5 to 1.6 V in an aqueous solution buffered at pH 8, the cationic dye MD+

yielded an electroactive polymer film on GC. Fig. 1 exhibits the cyclic voltammogram of the chemically

modified electrode, measured in a 0.2 M H2SO4 solution. The electropolymerized dye revealed a pair of

redox waves at a half-wave potential E1/2 of 0.18 V, which was 0.05 V more negative than the E1/2 of the

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monomeric dye held in solution. The inset in Fig.1 shows the cathodic peak current Ipc plotted against the

potential-cycle number at electropolymerization. The surface coverage of the polymer increased

proportionally with the cycle number up to about 80 cycles, with a polymer-growth rate of 5 × 10-11 mol

cm-2 per cycle. As the potential cycle was repeated more than 100 times, however, Θ gradually

approached a limiting value of 6 × 10-9 mol cm-2. This result suggests a low rate of electron transport

through the film, resulting from redox conduction.

Responding to the solution pH, the polymer was protonated and deprotonated. When the pH

increased from one to seven, the E1/2 of the polymer shifted to more negative potentials at a rate of 50

mV per unit pH (Fig.2). This change was similar to that (−57 mV/pH) observed for the monomeric dye

incorporated into Nafion, a perfluorosulfonated cation-exchange polymer. The latter rate is equal to that

expected for the two-electron, two-proton reduction of MD+ to the protonated leucoMeldola’s blue

MDH2+ (Scheme 1).

MD+ + 2e + 2H+ → MDH2+ (2)

This observation implies that the polymer includes the same electroactive site as the monomer.

According to the published acid dissociation constant Ka for amino-substituted azines38, 39), their

oxidized forms are unprotonated at pH values above three: their basicity is weaker than that of aniline. In

the present study, the pKa of the conjugate acid of MD+ was estimated to be less than 1 from the pH

dependence of the absorbance at an absorption peak wavelength of 569 nm, where MD+ showed a molar

absorption coefficient of 1.7 × 104 M-1cm-1.

MDH2+ → MD+ + H+ pKa < 1 (3)

The molar absorption coefficient decreased when MD+ was protonated in acid solutions. On the other

hand, reduction of the azine rings generally makes their amino substituents more basic, thus inducing

protonation of the amino groups31, 38, 40, 41). For example, protonated leucomethylene blue with a chemical

structure similar to MDH2+ has been reported to have a pKa of 5.623). From these observations, one can

safely state that the reduced Meldola’s blue is protonated at weakly acidic pH values.

Previous studies27) on the chemical structures of electropolymerized azines suggest that the amino

nitrogen atom of one azine molecule link to a heterocyclic carbon atom of another such molecule

(so-called head-to-tail bonding). Although our ATR-IR spectroscopic experiment failed to distinguish

between the dimethylamino and ring-bridging amino groups, for two reasons we assume that PMD

5

includes the primary structure of phenoxazine (possibly with the tertiary amino group being

demethylated27)). First, the polymer had voltammetric properties like those of the dye (in E1/2 and its pH

dependence). Second, visible absorption band of the polymer disappeared upon its reduction, indicating

the conversion of phenoxazine rings to unconjugated 10-hydro-phenoxazine rings. When the polymer is

doubly reduced, thus, two protons will enter the polymer film in the pH range from 1 to 7. Being

localized within the phenoxazine ring, its π-electrons are incapable of moving through the amino group

that bridges the two rings. The electroactive properties of PMD can result from electron self-exchange

between the isolated redox-active sites and from a high ionic conductivity of the polycationic material37).

A slow rate of the electron hopping causes unsuccessful electrodeposition of PMD films thicker than

10-8 mol cm-2. The high ionic conductivity is ascribed to a high concentration of charge-compensating

counterions in the film. If the polymer film of Θ = 5.5 × 10-9 mol cm-2 has a thickness of 17 nm, we can

estimate the electroactive-site concentration at 3.2 M, neglecting the film swelling caused by water.

At the cycle number not exceeding 100, the electroactive polymer film indicated a log(Ipc) versus

log(v) plot with a slope of 0.90 ± 0.08 (thin-film behavior) in the 2 to 200 mV s-1 range (the inset in

Fig.2). In this study, hence, the PMD-modified electrode will be treated as a diffusionless system, that is,

analyzed by the voltammetric theory for surface-confined species42). At sufficiently fast scan rates for the

peak-to-peak separation ∆Ep > 0.2/ n V to hold, the cathodic peak potential Epc is given by

Epc = E °′ − (RT /αnF )ln(αnFv /RTks) (4)

where E °′ is the formal potential, α is the transfer coefficient, ks is the rate constant of a surface reaction

(in s-1), v is the scan rate, n is the number of electrons transferred in the heterogeneous reaction, and all

other symbols have their usual meanings. If the condition for irreversible surface redox reactions is

experimentally satisfied, Eq.(4) enables one to determine ks from a variation in Epc with v. Fig. 3 shows

the potential difference (Epc − E °′) plotted against the logarithm of v for a 1.3 nmol cm-2 film in 0.2 M

Na2SO4 solutions. The pH 0.7 solution yielded a linear (Epc − E °′) versus log(v) plot at higher scan rates

than the pH 7 one did. The values of ks and α were determined from the slope of the straight line and

from its intercept with the horizontal line of (Epc − E °′) = 0. Table 1 lists the results obtained in an acidic

and neutral solutions. When the solution pH increased from 0.7 to 7, ks decreased to about one twenty,

with α remaining near 0.5.

As can be seen from Eq.(2), the redox reaction of PMD has the pseudo-first-order rate constant

including the proton concentration. Quick responses of Ipc and ∆Ep to the solution pH, examined in the

pH range from 1 to 7, suggest rapid penetration of protons through the film. They move within polyazine

6

films through a Grotthus-type process, via water and/or amine sites in the polymers43). The

potential-independent concentration of the supporting electrolyte can thus be distributed uniformly

throughout the PMD film. Consequently, acidic solutions result in a high proton concentration in the film,

increasing ks. On the other hand, protonation of the polymer (causing a rise in its positive charge) raises

the degree of swelling of the polymer film. This variation may increase the distance between

redox-active sites, thus reducing ks; however, this expectation is contrary to the experimental result. A

high ks value obtained in acidic solutions results from an increased proton concentration in the film.

3.2. Voltammetric characteristics of DA oxidation

Electrochemical properties of DA vary with the solution pH, because its oxidation corresponds to

the change of 1,2-benzenediol into 1,2-benzoquinone. Fig. 4 illustrates typical cyclic voltammograms

measured at a bare GC electrode in 1 mM DA solutions. At pH < 6, the electroactive species showed an

oxidation and rereduction waves having a ∆Ep of 0.3 to 0.4 V at v = 200 mV s-1, smaller than that

observed at platinum electrodes. Because the conjugate acid of DA has a pKa of 8.924), the pair of redox

waves is assigned to the two-electron, two-proton conversion of protonated DA into the quinone form

DAQ (Eq.(5), Scheme 2). When the pH was raised to 8, DA showed two additional waves: an anodic

peak at 0.6 V and a cathodic at -0.3 V. The positive potential scan to 0.7 V greatly lowered the

DA-regenerating peak appearing upon the reversed negative scan, indicating that the second

electrooxidation is followed by a chemical reaction. According to previous studies4, 11, 44), unprotonated

DAQ is cyclized at the amine side chain into 5,6-dihydroxyindoline (Eq.(6), Scheme 2), which is further

oxidized to indoline-5,6-quinone as shown by Eq.(7) (Scheme 2) 4, 11). In the present study, the repetitive

potential cycle between -0.3 and 0.7 V depressed voltammetric activity of the electrode gradually,

because of the deposition of insoluble products. An electrode-poisoning effect of DA oxidation in neutral

solutions has suggested that indoline-5, 6-quinone polymerizes to melanin-like compounds on electrode

surfaces and thus impedes interfacial charge transfer45). The intramolecular cyclization made the

following electrocatalytic oxidation of DA limit to pH values not more than 7.

It should be noted that the anodic peak potential Epa of DA moved to more positive values with

increasing its concentration (cb). Fig. 5 shows Epa versus cb plots measured at a bare and PMD-modified

electrodes in 0.2 M Na2SO4 solutions (pH 5). When cb rose from 0.1 to 2 mM, Epa changed from 0.30 to

0.42 V at v = 200 mV s-1 for the bare GC electrode, whereas it shifted gradually to less positive potentials

for the chemically modified one. At the bare GC electrode, a similar change was observed for linear

sweep voltammograms measured with a rotating disk electrode. The potential corresponding to a half of

7

the diffusion-limited current IL moved to more positive values with an increase in cb. In addition, this

potential shifted similarly when the rotation speed ω rose from 1.67 to 66.7 Hz (Fig.10 (a)). These

changes in potential noticed at the bare GC electrode can arise from the following reactions: (1)

interfacial charge transfer accompanied with chemical reactions, (2) the association of DA molecules at

higher concentrations to form aggregates, or (3) the adsorption of DA/DAQ on the GC surface.

Although examining the changes in Ipa/v1/2, Epa, and Ipa/Ipc with v, we failed to characterize the

mechanism of electrode reaction (5) aptly.

Because the third possibility is discussed in Section 3.3, let us consider the second possibility from

the limiting current measured at the rotating disk electrode. For the pseudo-first-order redox reaction, this

current increases proportionally with the square root of ω (units: Hz) 46):

IL = 1.554nFSD2/3η -1/6cbω1/2 (8)

where n denotes the total number of electrons transferred in the redox reaction of DA, D the diffusion

coefficient of the species in an aqueous solution, and η its kinematic viscosity. A proportional

relationship between IL and ω1/2 allows one to evaluate D from the slope. As given in Fig. 6, all log IL

versus log ω plots measured in the concentration range from 0.1 to 4 mM show linear lines with a slope

of 0.49 ± 0.06. When plotted against cb, the slopes of IL versus ω1/2 plots increase proportionally with the

concentration (the inset in Fig. 6); thus, oxidation of DA obeys good first-order kinetics. Equation (8)

gives a constant D of 6.5 × 10-6 cm2s-1, which agrees with the value reported by Chen and Cha7). These

results indicate that DA does not form any aggregates in the concentration range examined.

3.3. Electrocatalytic oxidation of DA at PMD-modified electrodes

3.3.1. Influence of PMD films on the kinetic parameters of DA oxidation

The redox polymer heightened the anodic peak current Ipa of DA measured by cyclic voltammetry.

Fig. 7 shows typical cyclic voltammograms observed at a polyazine-modified and bare electrodes in a 1

mM DA solution containing 0.2 M Na2SO4. Poly(methylene blue) hardly enhanced the current of DA

oxidation at any pH values, shifting the current peak to more positive potentials (Fig.7 a). On the other

hand, PMD raised Ipa of DA by a factor of 5 to 6 at a cb of 1 mM (Fig.7 b). However, its effect depended

on the DA concentration. At the PMD-modified electrode, Ipa of DA increased rapidly as cb rose from 0

to 0.2 mM, above which it approached a limiting value. At the bare GC electrode, Ipa of DA increased

linearly with cb up to 5 mM, indicating the first-order dependence on the concentration. As a result, the

8

ratio of the catalyzed to uncatalyzed peak currents dropped from 10 to 1 as cb rose from 0.1 to 5 mM (the

inset in Fig. 7 b).

The polymer-enhanced current will allow one to detect DA at concentrations of the order of 5 µM

by differential pulse voltammetry, which excludes capacitive currents effectively and gives peak-shaped

polarization curves. Fig. 8 shows both the differential pulse voltammograms and calibration plot for the

electrocatalytic oxidation occurring in dilute DA solutions. On renewed polymer film surfaces, Ipa had an

standard error of about ± 60% at 2 µM DA and not less than ± 100% at 1 µM. Hence, the

PMD-modified electrode enables one to determine DA quantitatively in the concentration range of 5 to

20 µM by the pulse technique, with a detection limit of 3 µM.

Rotating disk electrode voltammetry gives the well-defined solution hydrodynamics and thus

involves the simplest theoretical derivations for the kinetic analysis of electrocatalytic reactions. In this

stationary technique, we evaluated the rate constant of DA oxidation by using a PMD film with a Θ of

(4.8 - 6.1) × 10-9 mol cm−2 (showing thin-film behavior) to ensure fast charge transport through the

polymer. For mixed kinetic-diffusion control where the current I depends on the electrode potential E, I

is given by 47)

I = (1 − I /IL)Ik (9)

where Ik is the kinetic current corresponding to the rate of interfacial charge transfer. Because the DA

concentration at the electrode surface is zero at much more positive potentials than E °′, I is determined

by mass transport to the surface. From Eqs.(8) and (9), one obtains

1 /I = 1 /Ik + (1.554nFSD 2/3η -1/6cbω1/2) -1 (10)

The plot of I -1 against ω -1/2 thus enables one to evaluate Ik from the intercept at ω = ∞. At anodic

potentials where the reverse process is ignored, Ik can be written as

Ik = nFSk°cb exp[αnaF(E − E °′)/RT] (11)

where k° is the apparent standard heterogeneous rate constant including the proton concentration and na

is the number of electrons involved in the rate-determining step. If the logarithm of Ik increases linearly

with (E − E °′), then one can estimate k° and αna.

The rotating-disk voltammetric results obtained at the bare GC electrode indicated that Eq.(5)

9

comprised plural steps where one electron was transferred in the rate-determining step. Fig. 9 (a) shows

the I-1 versus ω -1/2 plots for the bare GC electrode in a 0.2 mM DA solution (pH 6). At less positive

potentials, a linear relationship was observed only in the higher ω range, as predicted from the

voltammograms in Fig. 10 (a). The plots measured at more positive potentials, on the other hand,

indicated a well-behaved linear relationship. The logarithm of Ik determined from the intercept at ω = ∞

is plotted against E in the inset in Fig. 9 (a). At much more positive potentials than E °′, the plots deviated

from a straight line because I approached IL. For the 0.2 - 2 mM DA solutions buffered at pH 6, we

estimated k° at (3.9 ± 0.7) × 10-4 cm s-1 and αna at 0.38 ± 0.08 from Eq.(11). The latter value suggests

that one electron be transferred in the rate-determining step. If assuming na = 2, then we obtain an α of

0.19 which is uncommon value48). A similar αna value of 0.40 was cyclic-voltammetrically obtained

from a linear change in Epa with the logarithm of v in the range from 0.2 to 10 V s-1. The result just

mentioned is consistent with the suggestion that electrochemical oxidation of DA to DAQ proceeds via

two consecutive one-electron transfer processes with the intermediate semiquinone4).

Voltammetric sensing based on polymer-modified electrodes depends largely on how rapidly the

polymer films communicate the electric signals stemming from target recognition to the underlying

electrodes. Fig. 10 (b) shows the linear sweep voltammograms of DA measured at a rotating

PMD-modified electrode in a 0.2 mM DA solution (pH 6). The polymer film shifted the E1/2 of DA to

less positive values (compare with the voltammograms in Fig.5), with IL remaining unchanged.

Moreover, the current rose to the limiting value in a narrow potential range compared with that for the

bare GC electrode. Fig. 9 (b) presents I -1 versus ω -1/2 plots measured at various potentials. The

polymer-coated electrode established a linear relationship between I -1 and ω -1/2 in the wider ω range than

the bare one did. The logarithm of Ik determined by extrapolation of the straight line is plotted as a

function of E in the inset in Fig. 9 (b). For a 0.2 mM DA solution (pH 6), we estimated k° at 5.0 × 10-3

cm s-1 and αna at 0.98 (which yields α = 0.49 for na = 2) from the straight line shown in the inset. This k°

value is an order of magnitude higher than that measured at the bare GC electrode. In addition, the

polymer film increased the dependence of Ik on E probably by enabling a two-electron transfer step. The

apparent standard heterogeneous rate constant is plotted against cb in Fig. 11. It decreased as cb was

raised to more than 0.5 mM, corresponding with the limiting Ipa observed at high concentrations.

Because both the redox reactions of DA and PMD include protons, the polymer film can change

the dependence of k° on the solution pH. In practice, further oxidation of DA to occur in basic solutions

made the reliable determination of k° limit to aqueous media more acidic than pH 7. Table 2

summarizes the results of the rotating disk electrode voltammetry performed at three different pH values.

Whereas the rate constant obtained at the bare GC electrode remains almost unchanged, that measured

10

at the polymer-modified electrode tends to decrease as the solution is acidified.

3.3.2. Interpretation of the electrocatalytic effect

Because the E °′ of the DA/DAQ couple (0.16 V at pH 7) is more positive than that of the

MD+/MDH couple (-0.19 V at pH 7), electron cross-exchange (12) is thermodynamically unfavorable,

MD+ + DA → MDH + DAQ (12)

that is, its standard Gibbs energy change ∆G ° = 2F [E °′(DA/DAQ) − E °′(MD+/MDH)] > 0. Hence, the

electrocatalytic oxidation of DA at PMD-modified electrodes should be attributed to other causes than

the cross-exchange reaction.

It should be noted that the polyelectrolyte film extracted DA from aqueous electrolyte solutions.

When having been soaked in 0.5 mM DA solutions for 1 hr, the polymer-modified electrode indicated

the anodic wave (at about 0.4 V at pH 5) due to the DA incorporated into the film. The surface

concentration Γ of DA calculated from the area under the wave was on the order of 4 × 10-10 mol cm-2

for a PMD film having a surface coverage of 5.4 × 10-9 mol cm-2. Assuming that the film is 17 nm thick,

we can estimate the DA concentration in the film at 2.4 × 10-4 mol cm-3 from the Γ value. When the

loaded film was soaked in a pure electrolyte solution, however, Γ decreased at a rate of 6−7% per minute.

This decrease indicates that the DA incorporated was gradually released from the polymer during the

slow potential scan. Thus, Γ values measured by this method were underestimated.

The surface concentration at equilibrium was determined by potential-step chronocoulometry,

assuming that the film has as small thickness as submonolayers. Voltammetric charge Q versus time t

curves measured at long times and a large-amplitude potential step are expressed by the Cottrell

equation43):

Q = Qa + 2nFScb(Dt /π)1/2 (13)

where Qa denotes the charge for oxidizing the electroactive species adsorbed on the electrode surface.

Fig. 12 shows typical Q versus t1/2 plots for a potential step from 0.25 to 0.65 V at pH 4. For bare

electrodes (the inset), the charge corrected for the residual charge at cb = 0 varied linearly with t1/2 in the

time range from 10 ms to 1 s, as expected for a diffusion-limited process. The straight line extrapolated

to zero time allows us to estimate both D from the slope and Γ from the intercept Qa = nFSΓ. The

diffusion coefficient estimated from the slope corresponded to a value of (5.8 ± 1.3) × 10-6 cm2s-1,

11

agreeing closely with that determined by the rotating disk voltammetry. Fig. 13 shows Γ plotted as a

function of cb. The chronocoulometric result indicated that DA is adsorbed onto the GC surface up to

about 4 × 10-11 mol cm-2. This interaction can be related to the ∆Ep values smaller than those measured at

platinum electrodes. The polymer-modified electrodes, on the other hand, showed linear Q versus t1/2

plots except for the early stages of the measurement (< 20 ms), with the intercept differing greatly from

zero. The deviation for a longer period than that observed at the bare electrode can result from mass

transport in the film. The surface concentration calculated from the intercept is plotted against cb in Fig.

13. It increases consistently up to 0.5 mM and then approaches to a limiting value. Hence, it is

reasonable to conclude that the decrease in k° with cb results from a limited DA concentration in the film.

The distribution constant of DA between the film and solution was roughly estimated at 3000 from the

slope of the linear segment at low concentrations in Fig. 13. The limiting surface concentration measured

for the modified electrode is an order of magnitude higher than that obtained for the bare electrode.

Consequently, the increases in k° and αna due to the surface modification can be ascribed to the binding

interaction of DA with the polymer film. It is noteworthy that this interaction accelerates the two-electron

transfer step.

To clarify a binding interaction between DA and the polymer, we discuss whether Γ is affected by

the ionic charge of the polymer or not. The surface concentration in the polymer film equilibrated with a

0.5 mM DA solution is plotted against its pH in the inset in Fig.13. It decreased as the pH was varied

from 5 to 0.7. When protonated in strongly acidic solutions, the polymer weakened the binding

interaction with the protonated DA probably because of electrostatic repulsion between the two

positively charged species. Thus, the chemical interaction of DA with PMD plays a key role in the

incorporation of the neurotransmitter with the polyelectrolyte film.

4. Conclusions

The present study leads to the following conclusions.

1. Whereas poly(methylene blue) hardly enhances the current of DA oxidation, PMD raises Ipa by a

factor of 10 at low cb values. This electrocatalytic effect of PMD enables one to detect DA at

concentrations of the order of 5 µM.

2. Because the rate of electron hopping in PMD is slower than that in poly(methylene blue), the

electrocatalytic activity is independent of electron self-exchange between neighboring redox-active sites.

3. The PMD film on electrode surfaces increases the k° of DA oxidation more than 10-fold and promotes

the two-electron transfer step, thus raising the dependence of Ik on the potential.

12

4. The increases in k° and αna due to the electrode modification arise from the incorporation of DA with

the polymer. Its protonation decreases the adsorption of DA to the film. A full understanding of the

mechanism of the incorporation requires spectroscopic studies of a chemical interaction between DA

and the polyelectrolyte.

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14

Figure Captions

Table 1 Rate constants of the surface redox reaction of PMD

Table 2 Rate constants of the oxidation of 0.2 mM DA at different pH values

Structure 1 Azines used in the present work.

Scheme 1 Redox reaction of Meldola’s blue.

Scheme 2 Possible scheme for oxidation of DA.

Fig. 1 Cyclic voltammogram measured at a PMD-modified electrode in a 0.2 M H2SO4 solution. Θ =

5.1 × 10-9 mol cm-2 (the potential-cycle number at electropolymerization = 100), v = 10 mV/s. The inset

shows an increase in Ipc with the potential-cycle number at electropolymerization. v = 5 mV/s.

Fig. 2 Half-wave potential plotted against the solution pH (the straight line shown has a slope of -50

mV/pH). The inset shows a log-log plot of Ipc versus v, with a slope of 0.91. Θ = 5.7 × 10-9 mol cm-2, a

0.2 M H2SO4 solution.

Fig. 3 Changes in Epc with the logarithm of v for a pH 7 (○) and 0.7 (□) solutions. Θ = 1.3 × 10-9 mol cm-2.

Fig. 4 Typical cyclic voltammograms measured at a bare electrode in 1 mM DA solutions buffered at

pH 5 () and 8 (-----). v = 100 mV/s.

Fig. 5 Changes in Epa with cb observed at a bare (○) and a PMD-modified (◊) electrodes in 0.2 M Na2SO4 solutions (pH 5). v = 200 mV/s.

Fig. 6 Log-log plots of IL versus ω measured at a bare electrode in the pH 6 solutions of different DA

concentrations. The inset shows a linear relationship between cb and the slope of IL-ω1/2 plots.

Fig. 7 Typical cyclic voltammograms observed at a polyazine-modified (solid line) and bare GC

(broken line) electrodes in a 1 mM DA solution containing 0.2 M Na2SO4. v = 100 mV/s. (a)

15

poly(methylene blue) (PMT); (b) PMD. The inset in Fig. 7b shows a rapid decrease in the catalytic

current with cb.

Fig. 8 Differential pulse voltammograms at progressively higher concentrations of DA and calibration

plot for the electrocatalytic oxidation in dilute DA solutions.

Fig. 9 Plots of I -1 against ω -1/2 measured at different potentials for (a) the bare GC electrode and (b)

the PMD-modified one in a 0.2 mM DA solution (pH 6). Measured potentials: (a) (○) 0.37 V, (•) 0.40,

(∆) 0.45, (□) 0.50, and (◊) 0.60; (b) (○) 0.25 V, (•) 0.275, (∆) 0.30, (□) 0.325, and (◊) 0.35. Insets of (a)

and (b) show the logarithm of Ik plotted against E respectively.

Fig. 10 Linear sweep voltammograms of a 0.2 mM DA solution (pH 6) measured at (a) a bare GC

electrode and (b) a PMD-modified one rotating at various rotation speeds/Hz.

Fig. 11 Standard heterogeneous rate constant plotted against cb. pH 6.

Fig. 12 Typical potential step Q versus t1/2 plots measured at the PMD-modified electrode in the pH 4

solutions of different DA concentrations. The inset shows Q versus t1/2 plots measured at the bare GC

electrode.

Fig. 13 Plots of Γ against cb measured at pH = 4. (◊) bare GC electrode and (○) PMD-modified one.

The inset shows the relationship between the solution pH and the Γ of DA adsorbed from 0.5 mM

solutions.

16

Table 1. Rate constants of the surface redox reaction of PMD

E1/2 /V vc /mV s-1 ks /s-1 α pH 0.7 0.171 320 13.9 0.55

7 -0.141 17 0.62 0.48

Takahiro YAMAGUCHI*, Teruhisa KOMURA, Sayomi HAYASHI, Manabu ASANO, Guang Yao NIU, Kohshin TAKAHASHI

17

Table 2. Rate constants of the oxidation of 0.2 mM DA at different pH values

bare GC PMD-modified electrode

αna k°/(cm s-1) αna k°/(cm s-1)

pH 6 0.38 3.9 × 10-4 0.98 5.0 × 10-3

pH 4 0.44 4.2 × 10-4 0.94 4.5 × 10-3

pH 2 0.46 4.6 × 10-4 0.86 2.1 × 10-3


18

S

N

(H3C)2N N(CH3)2

+O

N

N(CH3)2

+

Meldola’s blue Methylene blue

Structure 1 Takahiro YAMAGUCHI*, Teruhisa KOMURA, Sayomi HAYASHI, Manabu ASANO, Guang Yao NIU, Kohshin TAKAHASHI

19

O

N

N(CH3)2

H

O

N

N(CH3)2

+2e + 2H

++

H

Scheme 1 Takahiro YAMAGUCHI*, Teruhisa KOMURA, Sayomi HAYASHI, Manabu ASANO, Guang Yao NIU, Kohshin TAKAHASHI

20

N

OH

OH N

O

O(5)

H3

+

H3

+

-2e -2H+

N

O

O

OH

OH NH

(6)

H2

OH

OH NH

O

O NH

(7)-2e -2H+

Protonated DA Protonated DAQ

DAQ

N

OH

OH N

O

O(5)

H3

+

H3

+

-2e -2H+

N

O

O

OH

OH NH

(6)

H2

OH

OH NH

O

O NH

(7)-2e -2H+

Protonated DA Protonated DAQ

DAQ

Scheme 2 Takahiro YAMAGUCHI*, Teruhisa KOMURA, Sayomi HAYASHI, Manabu ASANO, Guang Yao NIU, Kohshin TAKAHASHI

21

0 0.5 1

/ V

10

0

/μA

E

I

0 100 2000

2

Cycle number

/μA

I p

Fig. 1


22

0 4 8-0.2

0

0.2

slope = -0.050

pH

1/2 /

VE

0 1 2

0

1

2

log( /mV s -1)

log(

p/μ

A)

slope = 0.91

v

I

Fig. 2


23

-2 0 2-0.2

-0.1

0

0.1

pH 7

pH 0.7

log( /V s -1)

( pc

- o ')

/V

v

EE

Fig. 3


24

-0.5 0 0.5-20

0

20

pH 5

pH 8

100 mV/s

1mM DA

/ V

/μA

E

I

Fig. 4


25

0 1 2

0.2

0.4

0.6

b / mM

pa/ V

PMD

bare

E

c0 1 2

0.2

0.4

0.6

b / mM

pa/ V

PMD

bare

E

c

Fig. 5 Takahiro YAMAGUCHI*, Teruhisa KOMURA, Sayomi HAYASHI, Manabu ASANO, Guang Yao NIU, Kohshin TAKAHASHI

26

0 1 2

2

4

log(ω/Hz)

log(

L/ μ

A)

0.1 mM

421

0.50.2

I

0 2 40

150

b / mM

Slop

e of

I Lvs

ω1/

2

plot

/(μ

As1/

2 )

bare, pH 6

c

Fig. 6

0 1 2

2

4

log(ω/Hz)

log(

L/ μ

A)

0.1 mM

421

0.50.2

I

0 2 40

150

b / mM

Slop

e of

I Lvs

ω1/

2

plot

/(μ

As1/

2 )

bare, pH 6

c

0 1 2

2

4

log(ω/Hz)

log(

L/ μ

A)

0.1 mM

421

0.50.2

I

0 2 40

150

b / mM

Slop

e of

I Lvs

ω1/

2

plot

/(μ

As1/

2 )

bare, pH 6

c

Fig. 6


27

-25

0

25

50

I / µ

A

-0.4 0 0.4 0.8E / V

-50

0

50

100

150

I / µ

A

-0.4 0 0.4 0.8E / V

PMT+DA

bare

PMT

PMD

bare

(a)

(b)

0

5

10

Rat

io o

f pe

sk c

urre

nts

0 2 4Cb / mM

Fig. 7

Takahiro YAMAGUCHI, Teruhisa KOMURA, Sayomi HAYASHI, Manabu ASANO, Guang Yao NIU, Kohshin TAKAHASHI

28

0 10 200

2

4

6

b /μM

Δ p

/μA

I

0.2 0.4 0.6 0.80

5

10

/ V

Δ /μ

A

20

5

0

100μM

E

I c

Fig. 8


29

0 0.4 0.80

50

100

150

(ω/Hz)-1/2

( /m

A)-1

I

0.2 0.3 0.4-4

-2

(E - Eo)/ V

ln(I

k/m

A)

(a)

0 0.4 0.80

50

100

150

(ω/Hz)-1/2

( /m

A)-1

I

0.05 0.1 0.15-3

-2

-1

(E - Eo)/V

ln(I

k/m

A)

Fig. 9

(b)

0 0.4 0.80

50

100

150

(ω/Hz)-1/2

( /m

A)-1

I

0.2 0.3 0.4-4

-2

(E - Eo)/ V

ln(I

k/m

A)

0 0.4 0.80

50

100

150

(ω/Hz)-1/2

( /m

A)-1

I

0.2 0.3 0.4-4

-2

(E - Eo)/ V

ln(I

k/m

A)

(a)

0 0.4 0.80

50

100

150

(ω/Hz)-1/2

( /m

A)-1

I

0.05 0.1 0.15-3

-2

-1

(E - Eo)/V

ln(I

k/m

A)

Fig. 9

(b)


30

0.2 0.4

0

20

40

60

1.67

3.33

8.33

16.7

33.3

66.7

/ V

/ μA

pH 6b = 0.2 mM

E

I

c

Fig. 10

(b)

0 0.4 0.8

0

30

60

/ V

/ μA

E

I(a)

1.67

3.33

8.33

16.7

33.3

66.7

pH 6

= 0.2 mMcb

0.2 0.4

0

20

40

60

1.67

3.33

8.33

16.7

33.3

66.7

/ V

/ μA

pH 6b = 0.2 mM

E

I

c

Fig. 10

(b)

0.2 0.4

0

20

40

60

1.67

3.33

8.33

16.7

33.3

66.7

/ V

/ μA

pH 6b = 0.2 mM

E

I

c

Fig. 10

(b)

0 0.4 0.8

0

30

60

/ V

/ μA

E

I(a)

1.67

3.33

8.33

16.7

33.3

66.7

pH 6

= 0.2 mMcb

0 0.4 0.8

0

30

60

/ V

/ μA

E

I(a)

1.67

3.33

8.33

16.7

33.3

66.7

pH 6

= 0.2 mMcb


31

0 1 20

2

4

6

b / mM

/10-3

cm s-1

o

pH 6

c

k

Fig. 11

0 1 20

2

4

6

b / mM

/10-3

cm s-1

o

pH 6

c

k

0 1 20

2

4

6

b / mM

/10-3

cm s-1

o

pH 6

c

k

Fig. 11Fig. 11


32

0 0.5 1

0.1

0.2

0.5

1

2 mM

/ μC

100

50

0

1/2 /s1/2t

Q

0 0.2 0.4 0.6

t1/2 /s1/2

/ μC

40

20

00.1

2 mM

0.5Q

Fig. 12

0 0.5 1

0.1

0.2

0.5

1

2 mM

/ μC

100

50

0

1/2 /s1/2t

Q

0 0.2 0.4 0.6

t1/2 /s1/2

/ μC

40

20

00.1

2 mM

0.5Q

0 0.5 1

0.1

0.2

0.5

1

2 mM

/ μC

100

50

0

1/2 /s1/2t

Q

0 0.2 0.4 0.6

t1/2 /s1/2

/ μC

40

20

00.1

2 mM

0.5Q


33

0 0.5 10

0.2

0.4

b /mM

/ nm

olcm

-2

PMD

bare

Γ

c

0 2 4 60

0.2

0.4

pH

/ nm

olcm

-2

PMD

bareΓ

Fig. 13

0 0.5 10

0.2

0.4

b /mM

/ nm

olcm

-2

PMD

bare

Γ

c

0 2 4 60

0.2

0.4

pH

/ nm

olcm

-2

PMD

bareΓ


34

Electrocatalytic activity of electropolymerized methylene blue … · 2019. 5. 10. · oxidation...

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Transcript of Electrocatalytic activity of electropolymerized methylene blue … · 2019. 5. 10. · oxidation...